Chapter 6 Chemical Composition

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Chapter 6 Chemical Composition 6.1 Counting by
Weighing Bean Lab We can count individual units
by weighing if we know the average mass of the
units. Instead of counting out 1000 jelly beans,
it is easier to (1.) find the mass of 1 bean (
5 g), (2.) multiply x 1000 (5000 g), and (3.)
measure out 5000 g of beans.
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6.2 Atomic Masses Counting Atoms by
Weighing Pennium Lab When we know the average
mass of the atoms of an element, we can calculate
the number of atoms in any given sample of that
element by weighing the sample. 1 amu 1.66
x10-24 g Using Table 6.1, p. 157, calculate the
mass, in amu of a sample of aluminum that
contains 75 atoms. 1 Al atom 26.98 amu mass
of 75 atoms 75 atoms x 26.98 amu 2024
amu 1 atom Calculate the number of sodium
atoms present in a sample that has a mass of
1172.49 amu, if 1 Na atom 22.99 amu. 1172.49
amu x 1 Na atom 51.00 Na atoms 22.99
amu
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6.3 The Mole 1 dozen 12 1 pair 2 1 score
20 1 mole 6.02 x 1023 objects (a counting
number, an exact number infinite of sig figs)
aka Avogadros Number (p. 160) A sample of an
element with a mass equal to that elements
average atomic mass expressed in grams contains 1
mole of atoms. Example see Table 6.2, p.
160 Conversions (p. 164) 5.00 x 1020 atoms
Cr Determine of moles Determine mass, in grams
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Mole Map Mass, in grams (use molar
mass) Mole (use
Avogadros ) (use 22.4 L) Number of
Atoms Volume, in Liters (of a gas at STP,
standard temperature, 0oC, and
standard pressure, 1 atm) (Ch. 13
Gases)
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6.4 Molar Mass Ionic compounds- calculate the
mass of 1 formula unit of the compound (compound
formula as written). For ex., NaCl molar mass
_________ CaCl2 molar mass _________ For
covalent compounds- calculate the mass of the
molecule (compound formula as written). For ex.,
(p. 166, ex. 6.5) SO2 molar mass 64.07
g C2H3Cl molar mass _________ CuSO4 5H2O
molar mass _________ Conversions mass from
moles, p. 168 Moles from mass, p. 168
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number of molecules (or formula units, if ionic)
from mass, p. 169
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6.5 Percent Composition of Compounds Percent (by
mass) for a given element mass of element x
100 mass of 1 mol of compound Example
MgCO3 molar mass ____________ Example
penicillin, p. 173
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• 6.6 Formulas of Compounds
• Empirical formula- lowest whole-number ratio of
  all elements in a compound formula. For ionic
  compounds, ALL formulas are empirical formulas.
• Molecular formula- for covalent compounds, the
  actual number of atoms of each type of element
  present in the compound.
• Ex. C6H12O6 is molecular formula CH2O is its
  empirical formula.

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6.7 Calculation of Empirical Formulas

• How to find
• Find actual of moles
• Find relative of moles (must be in whole
  numbers) divide actual of moles of each
  element by the smallest. Sometimes you must then
  multiply by some integer to get everything in
  whole numbers (integers) representing a ratio.
  These integers are your subscripts.
• Example p. 180
• Example p. 181
• Example p. 183 (using composition)(assume 100
  grams)

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6.8 Calculation of Molecular Formulas (all steps
are the same except for 1 additional step.) Check
the molar mass of the empirical formula if it is
the same as the molecular mass given in the
problem, then you are done. If not, you must find
by what factor is it different. Then multiply
the subscripts in the empirical formula by that
factor. Example p. 185